- Alan Chan
Since the Industrial revolution, there were great increases in emissions of sulfur dioxide and particulates, deteriorating air qualities more dominantly in industrial cities but also suburban areas and environments close by. Two of the very evident and dominant gases which cause much damage to our Earth include both sulfur dioxides and oxides of nitrogen which when reacted with water turn acidic through acid rain such that SO2(g) + H2O(l) –> H2SO3(aq) and 2NO2(g) + H2O(l) –> HNO2(aq) + HNO3(aq).
As shown, the reaction between the sulfur dioxides and oxides of nitrogen form acidic solutions when reacted with water as they release hydronium ions which indicate their acidic natures.
There are much of both (sulfur dioxide and oxides of nitrogen) produced naturally and also industrially.
Natural sources of sulfur dioxide contribute to 
of the total sulfur dioxide released to the atmosphere where it reacts with water and causes acid raid comes from activities of geothermal hot springs and volcanoes and the combustion of organic matter, eg bushfires and decomposition of organic matter.
Natural sources of the oxides of nitrogen, such as nitric oxide is lightning. These oxides of N2 are generated by lightning such that atmospheric oxygen and nitrogen gases combine to form nitric oxide:
O2(g) + N2(g) –> 2NO(g)
Then the nitric oxide slowly reacts with oxygen to form nitrogen dioxide:
2NO(g) + O2(g) –> 2NO2(g)
The above is the major natural source of nitrogen dioxide.
And finally, nitrous oxide is formed naturally by the action of certain bacteria on nitrogenous material in soils.
Main industrial origins of sulfur dioxide come from the combustion of fossil fuels (especially in power plants and motor vehicles). Smelting of sulphide ores during conversions of minerals to metals (such as lead, copper and zinc); incineration of garbage; petroleum refineries and industries using sulfur dioxide for production of sulfuric acid, production of paper, food processing and sewage treatment all contribute to the oxide of sulfur in the air.
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The main industrial origins of oxides of nitrogen include the large amounts of nitric oxide and nitrogen dioxide is combustion, both in stationary sources (power stations) and moving ones (motor vehicles). At high temperatures in combustion chambers, oxygen and nitrogen from air combine to form nitric oxide, and then nitric oxide is slowly converted to nitrogen dioxide. Releases of nitrous oxides to the atmosphere include the increased uses of nitrogenous fertiliser which provides more raw material for the bacteria.
In terms of concern for their release into the environment, there are many:
- Effects of sulfur dioxide and nitrogen oxides on human health ie; sulfur dioxide irritates the respiratory system and causes breathing difficulties at concentrations as low as 1ppm. Effects of sulfur dioxide are magnified if particulates are present also. Nitrogen dioxide irritates the respiratory tract and causes breathing discomfort at concentration levels of about 3-5ppm and greater that causes tissue damage.
- Effects of sulfur dioxide and nitrogen oxides on the environment -> formation of acid rain
H2O(l) + SO2 –> H2SO3(g) and 2NO2(g) + H2O(l) –> HNO2(aq) + HNO3(aq) which contributes to the increasing acidity of lakes (aquatic animals hence), damage to forests (such as pine forests in parts of Europe and North America), erosion of the marble and limestone of building surfaces and decorations and severe damage to vegetation especially around mine and smelter sites.
Even though this has happened very slowly over a long period of time, it is greatly concerning as these releases of gases will only increase as industries and mining in particular grows for the future. In the long term, these effects will become even more evident and greater and may be very difficult to restore or even slow down.
As mentioned before, the main contributor to high sulfur dioxide levels comes from the combustion of fossil fuels in power plants and motor vehicles. Coal and petroleum products contain sulfur, which combines with oxygen in air to form sulfur dioxide.
S(s) + O2(g) –> SO2(g)
Smelting of metal sulfides, eg copper sulphide, also produces sulfur dioxide.
CuS(s) + O2(g) –> SO2(g) + Cu(s)
Nitrogen undergoes combustion to form nitrogen monoxide and dinitrogen monoxide. Fossil fuels and biomass contain nitrogen (in proteins), so they also burn to produce these oxides.
N2(g) + O2(g) –> 2NO(g)
2N2(g) + O2(g) –> 2N2O(g)
N2(g) + 2O2(g) –> 2NO2(g)
Nitrogen monoxide burns to produce nitrogen dioxide.
2NO(g) + O2(g) –> 2NO2(g)
Although we may hear a lot about the effects of sulfur dioxides and oxides of nitrogen when reacted with water (acid rain), it is difficult to quantitatively state that oxides of sulfur and nitrogen have been increasing in the atmosphere because these oxides occur in relatively low concentrations, such as 0.01ppm, and the instruments used to measure these very low concentrations such as for SO2 have only been commercially available since the 1970s, so there is no reliable date for these gases before this time.
However, analysis of gases found in Antarctic ice core samples by the CSIRO and the Australian Antarctic Division showed that levels of N2O in the atmosphere has increased by about 10%.
Also, the increased burning of fossil fuels after the Industrial Revolution lead to a rise in oxides of sulfur, and evidence for this is the air quality of major industrial cities that deteriorated greatly.
Increase in acid rain, which is mainly caused by acidic oxides of nitrogen and sulfur dissolved in water, eg SO2(g) + H2O(l) –> H2SO3(aq) which is acidic and 2NO2(g) + H2O(l) –> HNO2(aq) + HNO3(aq), of which both are acidic.
This lead to the increasing damage to buildings, forests and aquatic organisms. Also, it was found that higher atmospheric concentrations of sulfur dioxide and nitrogen oxides in industrial areas than in non-industrial areas.
Although we do not have accurate measures of atmospheric oxides of sulfur and nitrogen taken over a long period of time, there is enough indirect evidence to conclude that significant increases in atmospheric concentrations of oxides of sulfur and nitrogen have indeed taken place, especially since industrial revolution
With the aforementioned much dangerous particulates, sulfur dioxides and oxides of nitrogen especially when reacted with water forming acid rain, it is very important to use indicators to determine if levels of pH in the environment are at damaging levels for organisms and growth of nature. A few of many every day uses of indicators include testing the pH (acidity/basicity) of water in aquariums and swimming pools, chemical wastes and soils.
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The testing of pH in the water of aquariums is extremely important as marine life are sensitive to changes in their water. If the water becomes too acidic or alkaline/basic, organisms including fish and plants may not be able to survive. A few drops of indicator can be placed in a sample of the water, or a pH paper already soaked in indicator can be used to measure the pH of the water.
As swimming pools are widely used by the public, the pH of the water also needs to be monitored regularly as it needs to be kept at almost neutral to avoid skin and eye irritations. Adding chlorine (hypochlorite ion) is one way to control acidity and stop algae from growing. The testing of pH will be similar to that of the testing of pH of water in aquarium.
Other than the public interactive environments, there are also chemical wastes which are produced industrially. The waste solutions from industries are tested before they are pumped into rivers or seas as they tend to be highly acidic. The pH of the wastes must be neutralised or they can be very harmful to the environment. The pH of chemical wastes is also measured by indicators, and substances are added to neutralise it.
Many plants only tolerate a narrow pH range, so the soil has to be tested regularly to ensure its survival. A way of testing the pH is to place a neutral white powder (such as barium sulfate or calcium sulfate) on top of moist soil, and then place a few drops of universal indicator on it. The powder then absorbs the moisture from the soil allowing the colour of the indicator to be clearly seen.
Other than sulfur dioxides and oxides of nitrogen when reacted with water which produces acid rain eg SO2(g) + H2O(l) –> H2SO3(aq) and 2NO2(g) + H2O(l) –> HNO2(aq) + HNO3(aq) and hence affecting society and environment, there are many other naturally occurring acids and bases. Some commonly known naturally occurring acids include hydrochloric acid, acetic acid, citric acid and ascorbic acid.
Hydrochloric acid HCl is produced by the glands in the lining of our stomachs to form an acidic environment for the breaking of complex food molecules by the enzymes.
Acetic acid CH3-COOH such that vinegar is about 4% solution of acetic acid and helps to preserve food. It is produced naturally by the bacterial action on alcohol in air.
Citric acid C6H8O7 is widespread in plant and animal tissue, especially in citrus fruit. It is also formed in our bodies during cellular respiration.
Ascorbic acid C6H8O6 also known as vitamin C is present in fresh fruits and vegetables. It is involved in many metabolic pathways and has an important role in healing, blood cell formation and tissue growth.
There are also many naturally occurring bases which may include ammonia, metallic oxides and carbonates.
Ammonia NH3 is present in the stale urine of humans and other animals. It is also formed through the anaerobic decay of organic matter.
Metallic Oxides such as iron(III) oxide, copper oxide and titanium (IV) oxides are insoluble and are solid bases found in minerals.
Carbonates such as calcium carbonate CaCO3 is found naturally as limestones.
As shown, there are many naturally occurring acids and this has been used to our advantage in homes such as using them as food additives. Acids are used as food additives to improve the taste and/or to preserve them. This is because many bacteria cannot survive in acidic conditions and if the acid used is weak enough and not harmful for human consumption, this will allow food to last over a period of time. Common acids for this use include acetic acid, citric acid and phosphoric acid.
Acetic acid(vinegar), phosphoric acid and citric acid is used to improve taste.
Propanoic acid is used as a preservative in bread.
Another advantage of having many acids is the use of them to make esters, which are “compounds formed when alkanoic acids react with alkanols, or more generally, when carboxylic acids combine with alcohols.” “Esters have pleasant, fruity odours and occur widely in nature as perfumes and flavouring agents”, hence its use as food additives, artificial fruit essences and in the manufacture of perfumes and cosmetics. As flavours, it is a combination of crude taste (sweet, salt, bitter) and odour, and it is these odours which contribute to flavours. It is easier to mass-produce these synthetic esters for use as flavours and perfumes in processed foods and cosmetics than to use naturally occurring ones and they represent little health hazards as “they contain only substances that occur in natural flavours”.
Overall, our environment is acidic as shown through the need to test the pH in our waters and soils, which was caused by sulfur dioxide, oxides of nitrogen and particulates released into the atmosphere where it reacts with water releasing hydronium ions through acid rain. These acids are very damaging to our society and environment in the long term as discussed through their lowering pH levels in waters, soils and corrosion of buildings. However, advantages of having acids were also discussed; being used to preserve foods, improve their tastes and using them as esters. The negative effects of acids may seem to outweigh the positive effects on a much wider scale but there are nonetheless both many advantages alongside the disadvantages which benefit our society and are essential to our everyday lives.
Bibliography
Excel HSC Chemistry – Jim Stamell – Reprinted 2012; P69-70
Conquering Chemistry HSC Course Fourth Edition – Roland Smith; P121-126,131-133.
Module 3: The Acidic Environment Theory Notes HSC Chemistry – Johnson for Irwin’s Atoms 2003
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